Transition Elements - Practice Questions & MCQ

In the extended form of the periodic table, the elements have been grouped into four blocks namely s, p, d and f-blocks. The elements belonging to groups 3to 12 are called d-block or transition elements. In these elements, the last electron enters (n-1) d-subshell. The configuration varies from (n-1)d¹ns² to (n-1)d¹⁰ns².  These are present between s-block and p-block elements. The properties of these elements are intermediate between the properties of s-block and p-block elements, i.e, d-block elements represent change or transition in properties from most electropositive s-block elements to least electropositive p-block elements. Therefore, these elements are called transition elements.
Earlier, the transition elements were regarded as those elements which possessed partially filled penultimate d-subshells in their ground state or in one of their commonly occurring oxidation states. This definition included coinage metals(Cu, Ag and Au) in the transition elements as their ions have partially filled (n-1)d-subshells although their atoms have filled(n-1) d-subshells in the ground state.
However, the above definition does not cover the elements of group 12, i.e, zinc metals(Zn, Cd and Hg) as these elements do not have partially filled (n-1) d-subshells either in the ground state or in ions.
However, zinc metals showing similarities in some of chemical properties with transition metals are also included in this block. These are considered as end members of the transition series in order to maintain a rational classification of elements.
Certain d-block elements are particularly important in living organisms. Iron, the transition element, is present in the largest quantity in the human body. The best known biological iron-containing compound is the protein haemoglobin, the red component of blood that is responsible for the transport of oxygen. Cobalt is the crucial element in vitamin B12, a compound that acts as a catalyst in the metabolism of carbohydrates, fats and proteins. Molybdenum and iron together with sulphur form the reactive portion of nitrogenase, a biological catalyst used by nitrogen-fixing organisms to convert atmospheric nitrogen into ammonia. Copper and zinc are important in other biological catalysts. Iron, zinc, copper, cobalt, nickel, manganese and molybdenum are known to be an essential component of enzymes. Vanadium and chromium are also essential for life. Some bad elements are also present in this block. For example, mercury is toxic and is a threat to the environment. 

The Lanthanide Contraction describes the atomic radius trend that the lanthanide series exhibit. The Lanthanide Contraction refers to the fact that the 5s and 5p orbitals penetrate the 4f sub-shell so the 4f orbital is not shielded from the increasing nuclear change, which causes the atomic radius of the atom to decrease. This decrease in size continues throughout the series.
The Lanthanide Contraction is the result of a poor shielding effect of the 4f electrons. The shielding effect is described as the phenomenon by which the inner-shell electrons shield the outer-shell electrons so they are not affected by nuclear charge. So when the shielding is not as good, this would mean that the positively charged nucleus has a greater attraction to the electrons, thus decreasing the atomic radius as the atomic number increases. The s orbital has the greatest shielding while f has the least and p and d in between the two with p being greater than d.

The Lanthanide Contraction can be seen by comparing the elements with f electrons and those without f electrons in the d block orbital. Pd and Pt are such elements. Pd has 4d electrons while Pt has 5d and 4f electrons. These 2 elements have roughly the same atomic radius. This is due to Lanthanide Contraction and shielding. While we would expect Pt to have a significantly larger radius because more electrons and protons are added, it does not because the 4f electrons are poor at shielding. When the shielding is not good there will be a greater nuclear charge, thus pulling the electrons in closer, resulting in a smaller than expected radius.

In general, ions of the same charge in a given series show progressive decrease in radius with increasing atomic number. This is because the new electron enters a d orbital each time the nuclear charge increases by unity. It may be recalled that the shielding effect of a d electron is not that effective, hence the net electrostatic attraction between the nuclear charge and the outermost electron increases and the ionic radius decreases. The same trend is observed in the atomic radii of a given series. However, the variation within a series is quite small. An interesting point emerges when atomic sizes of one series are compared with those of the corresponding elements in the other series. The curves in figure below, show an increase from the first (3d) to the second (4d) series of the elements but the radii of the third (5d) series are virtually the same as those of the corresponding members of the second series. This phenomenon is associated with the intervention of the 4f orbitals which must be filled before the 5d series of elements begin. The filling of 4f before 5d orbital results in a regular decrease in atomic radii called Lanthanoid contraction which essentially compensates for the expected increase in atomic size with increasing atomic number. The net result of the lanthanoid contraction is that the second and the third d series exhibit similar radii (e.g., Zr 160 pm, Hf 159 pm) and have very similar physical and chemical properties much more than that expected on the basis of usual family relationship.

The factor responsible for the lanthanoid contraction is somewhat similar to that observed in an ordinary transition series and is attributed to similar cause, i.e., the imperfect shielding of one electron by another in the same set of orbitals. However, the shielding of one 4f electron by another is less than that of one d electron by another, and as the nuclear charge increases along the series, there is fairly regular decrease in the size of the entire 4f norbitals.
The decrease in metallic radius coupled with increase in atomic mass results in a general increase in the density of these elements. Thus, from titanium (Z = 22) to copper (Z = 29) the significant increase in the density may be noted.

Metallic character

• All the transition elements or d-block elements are metals, since the number of electrons in the outermost shell is very small, i.e, either 1 or 2. They possess metallic properties such as:
  • High melting and boiling points.
  • Good conductors of heat and electricity.
  • Hard, malleable and ductile. Hg is an exception which is liquid and soft.
  • High density, metallic lustre and high enthalpies of atomization.
• They exhibit all the three types of structures, i.e, fcc, hcp and bcc.
• Both metallic and covalent bonding exists in transition metals. The metallic bonding is due to possession of one or two electrons in the outermost energy shell. They have low ionisation energies.

Enthalpy of atomisation
They have high enthalpies of atomisation which are shown in figure given below. The maxima at about the middle of each series indicate that one unpaired electron per d orbital is particularly favourable for strong interatomic interaction. In general, greater the number of valence electrons, stronger is the resultant bonding. Since the enthalpy of atomisation is an important factor in determining the standard electrode potential of a metal, metals with very high enthalpy of atomisation (i.e., very high boiling point) tend to be noble in their reactions (see later for electrode potentials).
Another generalisation that may be drawn from the given figure is that the metals of the second and third series have greater enthalpies of atomisation than the corresponding elements of the first series; this is an important factor in accounting for the occurrence of much more frequent metal – metal bonding in compounds of the heavy transition metals.

There is an increase in ionisation enthalpy along each series of the transition elements from left to right due to an increase in nuclear charge which accompanies the filling of the inner d orbitals. The variation in ionisation enthalpy along a series of transition elements is much less in comparison to the variation along a period of non-transition elements. The first ionisation enthalpy, in general, increases, but the magnitude of the increase in the second and third ionisation enthalpies for the successive elements, is much higher along a series.
The irregular trend in the first ionisation enthalpy of the metals of 3d series, though of little chemical significance, can be accounted for by considering that the removal of one electron alters the relative energies of 4s and 3d orbitals. You have learnt that when d-block elements form ions, ns electrons are lost before (n – 1) d electrons. As we move along the period in 3d series, we see that nuclear charge increases from scandium to zinc but electrons are added to the orbital of inner subshell, i.e., 3d orbitals. These 3d electrons shield the 4s electrons from the increasing nuclear charge somewhat more effectively than the outer shell electrons can shield one another. Therefore, the atomic radii decrease less rapidly. Thus, ionization energies increase only slightly along the 3d series. The doubly or more highly charged ions have dⁿ configurations with no 4s electrons. A general trend of increasing values of second ionisation enthalpy is expected as the effective nuclear charge increases because one d electron does not shield another electron from the influence of nuclear charge because d-orbitals differ in direction. However, the trend of steady increase in second and third ionisation enthalpy breaks for the formation of Mn2+ and Fe3+ respectively. In both the cases, ions have d⁵ configuration. Similar breaks occur at corresponding elements in the later transition series.
The three terms responsible for the value of ionisation enthalpy are attraction of each electron towards nucleus, repulsion between the electrons and the exchange energy. Exchange energy is responsible for the stabilisation of energy state. Exchange energy is approximately proportional to the total number of possible pairs of parallel spins in the degenerate orbitals. When several electrons occupy a set of degenerate orbitals, the lowest energy state corresponds to the maximum possible extent of single occupation of orbital and parallel spins (Hunds rule). The loss of exchange energy increases the stability. As the stability increases, the ionisation becomes more difficult. There is no loss of exchange energy at d⁶ configuration. Mn⁺ has 3d⁵4s¹ configuration and configuration of Cr⁺ is d⁵, therefore, ionisation enthalpy of Mn⁺ is lower than Cr⁺. In the same way, Fe²⁺ has d⁶ configuration and Mn²⁺ has 3d⁵ configuration. Hence, ionisation enthalpy of Fe²⁺ is lower than the Mn²⁺. In other words, we can say that the third ionisation enthalpy of Fe is lower than that of Mn.

One of the notable features of a transition elements is the great variety of oxidation states these may show in their compounds.
The elements which give the greatest number of oxidation states occur in or near the middle of the series. Manganese, for example, exhibits all the oxidation states from +2 to +7. The lesser number of oxidation states at the extreme ends stems from either too few electrons to lose or share (Sc, Ti) or too many d electrons (hence fewer orbitals available in which to share electrons with others) for higher valence (Cu, Zn). Thus, early in the series scandium(II) is virtually unknown and titanium (IV) is more stable than Ti(III) or Ti(II). At the other end, the only oxidation state of zinc is +2 (no d electrons are involved). The maximum oxidation states of reasonable stability correspond in value to the sum of the s and d electrons upto manganese (Ti^IV O₂, V^VO₂⁺, Cr^V1O₄^2–, Mn^VIIO₄^–) followed by a rather abrupt decrease in stability of higher oxidation states, so that the typical species to follow are Fe^II,III, Co^II,III, Ni^II, Cu^I,II, Zn^II.
The variability of oxidation states, a characteristic of transition elements, arises out of incomplete filling of d orbitals in such a way that their oxidation states differ from each other by unity, e.g., V^II, V^III,V^IV, V^V. This is in contrast with the variability of oxidation states of non transition elements where oxidation states normally differ by a unit of two.
An interesting feature in the variability of oxidation states of the d–block elements is noticed among the groups. Although in the p–block the lower oxidation states are favoured by the heavier members (due to inert pair effect), the opposite is true in the groups of d-block. For example, in group 6, Mo(VI) and W(VI) are found to be more stable than Cr(VI). Thus Cr(VI) in the form of dichromate in acidic medium is a strong oxidising agent, whereas MoO₃ and WO₃ are not.
Low oxidation states are found when a complex compound has ligands capable of π-acceptor character in addition to the σ-bonding. For example, in Ni(CO)₄ and Fe(CO)₅, the oxidation state of nickel and iron is zero.

When a magnetic field is applied to substances, mainly two types of magnetic behaviour are observed: diamagnetism and paramagnetism. Diamagnetic substances are repelled by the applied field while the paramagnetic substances are attracted. Substances which are attracted very strongly are said to be ferromagnetic. In fact, ferromagnetism is an extreme form of paramagnetism. Many of the transition metal ions are paramagnetic.
Paramagnetism arises from the presence of unpaired electrons, each such electron having a magnetic moment associated with its spin angular momentum and orbital angular momentum. For the compounds of the first series of transition metals, the contribution of the orbital angular momentum is effectively quenched and hence is of no significance. For these, the magnetic moment is determined by the number of unpaired electrons and is calculated by using the ‘spin-only’ formula, i.e.,

where n is the number of unpaired electrons and μ is the magnetic moment in units of Bohr magneton (BM). A single unpaired electron has a magnetic moment of 1.73 Bohr magnetons (BM).
The magnetic moment increases with the increasing number of unpaired electrons. Thus, the observed magnetic moment gives a useful indication about the number of unpaired electrons present in the atom, molecule or ion.